Problem 48 - Chemical Principles: The Quest for Insight

Question

Write Lewis structures, including typical contributions to the resonance structure (where appropriate, allow for the possibility of octer expansion), for (a) dihydrogen phosphate ion; (b) sulfite ion; (c) chlorate ion; (d) nitrate ion.

Step-by-Step Solution

Answer

Dihydrogen phosphate ion has one P atom bonded to 2 OH groups and 2 O atoms, sulfite ion includes S with 2 O atoms and 1 O with a negative charge (resonance), chlorate ion contains Cl bonded to 3 O atoms (resonance), and nitrate ion consists of N bonded to 3 O atoms with resonance structures involving the delocalization of the double bond.

Step 1: Dihydrogen Phosphate Ion (H2PO4-)

Begin by counting the valence electrons. Phosphorus (P) has 5, each hydrogen (H) has 1, and oxygen (O) has 6. Since there are two hydrogens, four oxygens, and one negative charge, the total is 5 + (1*2) + (6*4) + 1 = 32 valence electrons. Phosphorus is the central atom with oxygens surrounding it and hydrogens attached to two oxygens. Two of the oxygens have single bonds with hydrogen, and the other two form double bonds with phosphorus. All atoms satisfy the octet rule except for the hydrogens, which have a full valence shell with 2 electrons.

Step 2: Sulfite Ion (SO3^2-)

Count the valence electrons. Sulfur (S) has 6 and oxygen (O) has 6. Adding the extra 2 electrons for the charge, we have a total of 6 + (6*3) + 2 = 26 valence electrons. Sulfur is at the center, bonded to three oxygens. One double bond and two single bonds are initially formed, with the extra electrons placed on the oxygens to give them full octets. The negative charges are on the oxygen atoms with single bonds. Resonance structures can be created by moving the location of the double bond among the three oxygens.

Step 3: Chlorate Ion (ClO3-)

Count valence electrons. Chlorine (Cl) has 7 and oxygen (O) has 6. With the ion's charge, we have 7 + (6*3) + 1 = 26 valence electrons. Chlorine is the central atom, surrounded by three oxygens. One double bond and two single bonds to the oxygens are made, with the extra electrons placed on the oxygens to complete their octets. The negative charge is on the oxygen with a single bond. Resonance structures involve shifting the double bond between the different oxygen atoms, similar to the sulfite ion.

Step 4: Nitrate Ion (NO3-)

Calculate the valence electrons. Nitrogen (N) has 5 and oxygen (O) has 6. Considering the charge, there are 5 + (6*3) + 1 = 24 valence electrons. Nitrogen is the central atom, with three oxygens around it. One double bond and two single bonds are made, and the extra electrons are placed on the oxygens with single bonds to fill their octets. The single negative charge can be delocalized, leading to resonance structures with the double bond rotating among the three oxygens.

Key Concepts

Valence Electrons

Valence electrons are the electrons in an atom that are found in the outermost shell and are involved in forming bonds with other atoms. Each element has a characteristic number of valence electrons; for example, hydrogen (H) has 1, while oxygen (O) has 6, phosphorus (P) has 5, sulfur (S) has 6, chlorine (Cl) has 7, and nitrogen (N) has 5.

Understanding the count of valence electrons is crucial when drawing Lewis structures, as it determines how atoms bond and what molecular shape they will form. In the provided exercise, the correct calculation of valence electrons for each ion is the first step to drawing accurate Lewis structures. For instance, for the dihydrogen phosphate ion (H₂PO₄^-), the total count of 32 valence electrons reflects the sum of the valence electrons from phosphorus, oxygen, and hydrogen, plus the additional electron due to the negative charge.

The concept of valence electrons is closely tied to the octet rule, as most atoms aim to have eight electrons in their valence shell through bonding - the main exception being hydrogen, which only requires two electrons to achieve a full valence shell, analogous to the helium noble gas configuration.

Resonance Structures

Resonance structures are a way of describing bonding in certain molecules or ions by delocalization of electrons resulting in the formation of multiple valid Lewis structures. These structures are all valid representations that contribute to the real, intermediate structure of the molecule. It's like having different snapshots of the molecule, each showing alternative placements of electrons.

In the exercise, you can see how resonance structures come into play while considering ions like sulfite (SO₃^2-) and nitrate (NO₃^-). These ions have multiple ways to arrange the double bonds without altering the position of the atoms themselves, thus different resonance structures can be drawn. In short, these structures demonstrate the concept of delocalized electrons within a molecule, contributing to the chemical stability and reactivity of the molecule. When drawing resonance structures, one must ensure that the positions of the nuclei remain constant while the electrons are allowed to 'move' between equivalent bonding situations.

Octet Rule

The octet rule is a chemical rule of thumb that reflects the tendency of atoms to bond in such a way that they each have eight electrons in their valence shell, giving them the same electronic configuration as a noble gas. The rule applies to most of the main-group elements, with notable exceptions for hydrogen (which follows the 'duet rule'), helium, and elements in the third period or higher that can expand their octet due to available d-orbitals.

In our exercise, the octet rule plays a fundamental role in determining how atoms will bond in the various ions. Phosphorus in the dihydrogen phosphate ion, for example, follows the octet rule by forming double bonds with oxygen, allowing it to have a stable arrangement of eight electrons around it. However, in some cases, elements can go beyond the octet rule, as seen with sulfur in the sulfite ion or chlorine in the chlorate ion, which have an expanded octet due to the extra electrons contributed by the negative charge. It is essential to understand the octet rule when drawing Lewis structures, but also to recognize the exceptions, so that accurate valence configurations are offered for elements capable of octet expansion.